Nafeesa

OBJECTIVE

To determine the concentrations of two acids through acid-base titration.

OVERVIEW

An acid-base titration is a method used to quantitatively determine the concentration of an unknown acid or base (the analyte) by neutralizing it with a solution of known concentration (the titrant). The equivalence point can be found either by identifying the steepest point on the titration curve measured with a pH meter, or by the color change of a pH indicator.

In this experiment, two systems will be examined:

1. Strong acid with strong base. A hydrochloric acid (HCl) solution of unknown concentration will be titrated with a standard 0.100 M sodium hydroxide (NaOH) solution.
   Note: the unknown HCl concentration is close to that of human stomach acid.
2. Weak acid with strong base. An acetic acid (CH₃COOH) solution of unknown concentration will be titrated with a standard 0.100 M NaOH solution.
   Note: the unknown CH₃COOH concentration is about ten times more dilute than typical household vinegar.

THEORY

The Dissociation Constant of Water

Because of the auto-ionization of water, aqueous solutions always contain both hydronium ions and hydroxide ions:

2 H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)

The H⁺ ion is hydrated immediately, joining a second water molecule (the conjugate base, in Lewis terms) to form H₃O⁺, the hydronium ion.

Under the Brønsted–Lowry definition, an acid is any substance that can donate a proton (H⁺) to another species. In aqueous solution, a B-L acid is any substance that raises the H⁺(aq) concentration (by producing H₃O⁺(aq) ions). Likewise, a B-L base is a substance that can accept a proton and, as a result, lowers the H₃O⁺(aq) concentration while raising the OH⁻(aq) concentration.

In every aqueous solution, [H₃O⁺] and [OH⁻] are not independent. They are connected by:

K_w = [H₃O⁺][OH⁻] = 1 × 10⁻¹⁴

where square brackets denote concentrations in moles per litre (M). This relationship is strictly true only at 25 °C, though it remains approximately valid at temperatures close to room temperature.

The pH Concept

The range of [H₃O⁺] values is very broad, so a logarithmic scale is more convenient. pH (“power of hydrogen”) is defined as:

pH = −log[H⁺]

A pH of 7 is neutral (this includes pure water). Solutions with pH below 7 are acidic, and those above 7 are basic. From the relationship between [H₃O⁺] and [OH⁻], one can derive:

pH + pOH = 14

Titration of Acids and Bases

An acid-base titration is carried out to find the unknown concentration of an acid or base. At the equivalence point, for a monoprotic acid and a monohydroxy base:

M(standard) × V(standard) = M(unknown) × V(unknown)

where M is molarity and V is volume.

By definition, in every neutralization reaction the equivalence point is where the moles of H₃O⁺ equal the moles of OH⁻. For example, with NaOH and HCl:

NaOH(aq) + HCl(aq) → H₂O(l) + NaCl(aq)

Because NaOH, HCl and NaCl are all fully dissociated in aqueous solution, the net reaction is:

OH⁻(aq) + H₃O⁺(aq) → 2 H₂O(l)

To determine the concentration of the unknown, you must know precisely the number of moles of ions in the standard, together with the volumes of both the standard and the unknown. You also need some way of detecting when the equivalence point has been reached.

For the first requirement, a precisely known volume of standard is used. This volume (in litres) multiplied by the concentration gives the moles of ions used. This equals the moles of ions ultimately neutralized in the unknown, whether the unknown is a weak or strong acid.

Determining the equivalence point is more complex and depends on whether the unknown is weak or strong. (Note: a strong acid or base is generally chosen as the standard; only the unknown may be weak or strong.)

Strong Acid – Strong Base Titration

Standard NaOH is used to titrate unknown HCl:

NaOH(aq) + HCl(aq) → H₂O(l) + NaCl(aq)

At equivalence, all of the original acid has been consumed and no excess base is present. The only species present are H₂O and NaCl. Because NaCl is neutral (as is water), nothing pushes the pH into the acidic or basic range. In a titration of a strong acid with a strong base, the equivalence point is at pH = 7.

Weak Acid – Strong Base Titration

Standard NaOH is used to titrate an unknown amount of CH₃COOH:

CH₃COOH(aq) + NaOH(aq) → H₂O(l) + CH₃COONa(aq)

The species present at equivalence are H₂O and CH₃COONa. The pH at equivalence is not 7, because CH₃COONa acts as a B-L base, lowering the H₃O⁺ concentration while raising the OH⁻ concentration. The pH at equivalence therefore depends on the concentration of CH₃COONa.

Locating the Endpoint in an Acid-Base Titration

1. Using indicators. Acid-base indicators are complex organic compounds (themselves weak acids or bases) that change color as the pH changes. If you choose an indicator whose color-change pH lies on the vertical section of the titration curve, the endpoint shown by the indicator will match the equivalence point.

2. Using a pH meter. A pH meter has an electrode that produces an electrical signal related to the H⁺(aq) concentration of the solution. The signal is calibrated to give readings (analogue or digital) directly in pH units. The meter is placed in the analyte solution and records the pH as titrant is added. The endpoint is the point where the pH changes most rapidly with a small addition of titrant, i.e. the steepest point on the titration curve.

PRE-LABORATORY PREPARATION

Safety Data Sheets (SDS)

Refer to online SDS databases (e.g. www.google.ca, searching the chemical’s name plus “SDS”) for: sodium hydroxide, acetic acid, hydrochloric acid, phenolphthalein and methyl orange. Pay particular attention to “Section 2 – Hazard(s) identification” and “Section 7 – Handling and storage”.

PROCEDURE

Part A. Titration with a pH Meter

1. Check with your lab demonstrator whether the pH meter has been calibrated.
2. Obtain 70 mL of “NaOH Standard (0.100 M)” in a clean 100-mL beaker.
3. Rinse the buret with 10 mL of “NaOH Standard (0.100 M)”. Dispose of the rinse in your personal waste beaker.
4. Fill the buret with “NaOH Standard (0.100 M)”, with the top of the solution slightly above the 0.00 mL mark. Drain a small amount into your personal waste beaker to fill the buret tip. Adjust the solution to exactly 0.00 mL (bottom of the meniscus on the 0.00 mL line).
5. Obtain 20 mL of “HCl unknown” in a clean 50-mL beaker. Rinse the transfer pipet with 5 mL of “HCl unknown” and dispose of the rinse in the personal waste beaker.
6. Pipet 10.00 mL of “HCl unknown” into a clean 250-mL beaker. Dilute to roughly 100 mL with distilled water. Add a magnetic stirring bar and set the beaker on a magnetic stirrer.
7. Take the pH meter out of the KCl beaker. (Note: never let the pH meter dry out in air. Keep it in the KCl beaker between measurements and rinse it with distilled water before use.) Rinse it with distilled water and dispose of the rinse in your personal waste beaker.
8. Submerge the pH meter in the “HCl unknown”. Switch on the magnetic stirrer. Make sure the stirring is gentle, with no bubbling, and that the stir bar does not strike the pH meter.
9. Assemble the apparatus as shown in the figure, then have your lab demonstrator verify it.

Part B. Titration with a pH Indicator (Phenolphthalein)

1. Take about 20 mL of “HCl unknown” in a clean 50-mL beaker. Rinse the transfer pipet with about 5 mL of “HCl unknown” and pour the rinse into your personal waste beaker.
2. Pipet 10.00 mL of “HCl unknown” into a clean 250-mL Erlenmeyer flask. Dilute to roughly 100 mL with distilled water. Add 3 drops of phenolphthalein and swirl for 10 seconds.
3. Measure about 70 mL of “NaOH Standard (0.100 M)” into a clean 100-mL beaker.
4. Rinse the buret with about 10 mL of “NaOH Standard (0.100 M)” and pour the rinse into your personal waste beaker.
5. Fill the buret with “NaOH Standard (0.100 M)”, with the top of the solution slightly above the 0.00 mL mark. Let a small amount run into your personal waste beaker to fill the tip. Adjust to exactly 0.00 mL (bottom of the meniscus on the 0.00 mL line).
6. Arrange the apparatus as shown in the figure, then have your lab demonstrator verify it.
7. Read the buret (to two decimal places, e.g. 0.00 mL) and record it as “V initial”.
8. Open the stopcock to let about 6 mL of NaOH flow in. The solution in the flask will turn pink. Swirl for 5 seconds, and the pink color will fade.
9. Open the stopcock slowly and add NaOH drop by drop (about one drop every two seconds), swirling gently. A faint pink color will appear and fade quickly. Watch the buret reading closely and close the stopcock when it reaches 9 mL.
10. The titration is now very near the endpoint. Adjust the stopcock to let just one drop fall at a time, swirling for 3 seconds after each. The color will fade more slowly. If liquid remains on the buret tip, rinse it into the flask with deionized water from a wash bottle so all the NaOH ends up in the reaction mixture.
11. The endpoint is reached when the faint pink color persists for at least 30 seconds.
12. Read the buret (to two decimal places, e.g. 10.00 mL) and record it as “V final”. Calculate:
    V(NaOH standard) = V final − V initial
13. Determine the concentration of the “HCl unknown”:
    M(NaOH standard) × V(NaOH standard) = M(HCl unknown) × V(HCl unknown)
    where M(NaOH standard) = 0.100 mol/L and V(HCl unknown) = 10.00 mL, so:
    M(HCl unknown) = 0.100 mol/L × V(NaOH standard) ÷ 10.00 mL
    Example: if V(NaOH standard) = 10.00 − 0.00 = 10.00 mL, then
    M(HCl unknown) = 0.100 × 10.00 ÷ 10.00 = 0.100 mol/L
14. Repeat steps 1 to 13 using 10.00 mL of “Acetic Acid unknown” in place of “HCl unknown”.

Main Applications of Acid-Base Titration

Pharmaceuticals

Titration lets laboratories verify that a medicine contains the correct amount of its active ingredient, such as aspirin, ibuprofen, vitamin C or ephedrine hydrochloride. It is also used to calibrate the concentrations of lab reagents and to check raw drug mixtures for unwanted contaminants.

Food and Drink

Producers test acidity to judge freshness, flavor and how long a product will keep. Typical examples are lactic acid in fermented dairy, and citric or acetic acid in fruit juices and pickled foods. Winemakers and cheesemakers also track acidity as their products age.

Environment and Water Quality

Technicians use titration to find out how acidic or alkaline lake, river, waste and pool water is, so its pH can be corrected where needed. It also helps detect pollution and contaminated runoff that disturb natural water bodies.

Industrial Manufacturing

Manufacturers check the leftover alkalinity in soaps, detergents, cosmetics and dyes to confirm product quality. The method is also used to measure the strength of acids that feed chemical processes, including those used to make polymers.

Agriculture and Soil Science

Farmers and soil scientists measure soil acidity and its resistance to pH change (buffer capacity), which helps them create better conditions for crops.